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Chapter 16 Acid-Base Equilibria

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1 Chapter 16 Acid-Base Equilibria
AP CHEMISTRY Chapter 16 Acid-Base Equilibria Content from The Central Science, 9th ed David P. White, revised 2004 by CMB Chapter 16-Acids and Bases

2 Acids and Bases: A Brief Review
Acids: taste sour and cause dyes to change color. Bases: taste bitter and feel soapy. Arrhenius Definition: acids increase [H+] bases increase [OH-] in solution. Arrhenius: acid + base  salt + water. Arrhenius limits us to aqueous solution. Chapter 16-Acids and Bases

3 Brønsted-Lowry Acids and Bases
Hydronium: The H+ Ion in Water The H+(aq) ion is simply a proton with no electrons. Generally we use H+(aq) and H3O+(aq) interchangeably. Brønsted-Lowry: acid donates H+ and base accepts H+. Brønsted-Lowry base does not need to contain OH-. Consider HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq): HCl donates a proton to water. Therefore, HCl is an acid. H2O accepts a proton from HCl. Therefore, H2O is a base. Water can behave as either an acid or a base-its amphoteric. Amphoteric substances can behave as acids and bases. Example on next page Chapter 16-Acids and Bases

4 HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq):

5 The Autoionization of Water
The Ion Product of Water In pure water the following equilibrium is established at 25 C The above is called the autoionization of water. In neutral, pure water, [H3O+] = 1 x 10-7 M In neutral, pure water, [OH-] = 1 x 10-7 M Chapter 16-Acids and Bases

6 Chapter 16-Acids and Bases
The pH Scale “p” = -log In most solutions [H+(aq)] is quite small. We define In neutral water at 25 C, pH = pOH = 7.00. In acidic solutions, [H+] > 1.0  10-7, so pH < 7.00. In basic solutions, [H+] < 1.0  10-7, so pH > 7.00. The higher the pH, the lower the pOH, the more basic the solution. There are no theoretical limits on the values of pH or pOH. (e.g. pH of 2.0 M HCl is ) Logarithm algebra is needed here Chapter 16-Acids and Bases

7 The pH Scale, & Measuring pH
pH meter – measures conductivity as a function of [H3O+]. Precise, but calibration is often needed for accuracy Indicators – dyes that change their bonding & color as a function of [H3O+]. Less precise than pH meters, but quick and cheap Most indicators tend to be red in more acidic solutions. Chapter 16-Acids and Bases

8 Indicators & the pH Scale
Endpoint: the pH that produces a noticeable color change Transition Interval: the range of pH that they are effective.

9 Conjugate Acid-Base Pairs
Definition Conjugate Base = Whatever is left of the acid after the proton is donated. Conjugate Acid = Whatever remains of the base after it accepts a proton. Consider After HA (acid) loses its proton it is converted into A- (base). Therefore HA and A- are conjugate acid-base pairs. After H2O (base) gains a proton it is converted into H3O+ (acid). Therefore, H2O and H3O+ are conjugate acid-base pairs. Conjugate acid-base pairs differ by only one proton. Chapter 16-Acids and Bases

10 Identify Conjugate pairs:
Acid Base Conjugate Acid Conjugate Base Base Acid Conjugate Base Conjugate Acid Chapter 16-Acids and Bases

11 Relative Strengths of Acids and Bases
Strengths of Acids or Bases is determined by the extent to which a species dissociates, or ionizes in water. Chapter 16-Acids and Bases

12 Conjugate Acids and Base Strength
Rules: Relative Strengths of Acids and Bases The stronger the acid, the weaker the conjugate base. The stronger the base, the weaker the conjugate acid. A “strong” acid’s conjugate base (e.g. Cl-) has negligible acid-base properties. A “strong” base’s conjugate acid has negligible acid-base properties. Conjugates of weak acids are weak bases. Conjugates of weak bases are weak acids. Acid strength increases across a period and down a group. Base strength decreases across a period and down a group. Chapter 16-Acids and Bases

13 Strong Acids and Bases *MEMORIZE*
Six Strong Acids*: HClO4 HI HBr HCl H2SO4 HNO3 HNO3(aq)  H+(aq) + NO3-(aq) Strong Bases: NaOH KOH Ba(OH)2 etc-OH (those soluble) Strong bases need not have OH- (e.g.: H-(aq) + H2O(l)  H2(g) + OH-(aq) ) All others, consider to be “weak” (not completely dissociated) Strong acids & bases are strong electrolytes. “Weak” acids and bases = Ions are in equilibrium with unionized compounds. *Your book: HClO3 is also strong, but not on AP test. Chapter 16-Acids and Bases

14 Chapter 16-Acids and Bases
Weak Acids The of the acid-base on the AP Weak acids are only partially ionized in solution. Therefore, weak acids are in equilibrium: Ka is the acid dissociation constant. The larger the Ka the stronger the acid (i.e. the more ions are present at equilibrium relative to unionized molecules). If Ka >> 1, then the acid is completely ionized and the acid is a strong acid. Chapter 16-Acids and Bases

15 Chapter 16-Acids and Bases
Some Weak Acids: Chapter 16-Acids and Bases

16 Calculating pH or [H3O+] from Ka
Weak acids are simply equilibrium calculations. The pH gives the equilibrium concentration of H+. Using Ka, the concentration of H+ (and hence the pH) can be calculated: Write the balanced chemical equation clearly showing the equilibrium. Write the equilibrium expression. Find the value for Ka. Write down the initial and equilibrium concentrations for everything except pure water. We usually assume that the change in concentration of H+ is x. Substitute into the equilibrium constant expression and solve. Remember to turn x (at equilibrium) into pH if necessary. We can also solve for Ka from pH or [H3O+] in this manner. Chapter 16-Acids and Bases

17 Chapter 16-Acids and Bases
% Ionization Using %Ionization to Calculate pH or Ka Percent ionization is another method to assess acid strength. For the reaction The higher percent ionization, the stronger the acid. Percent ionization of a weak acid decreases as the molarity of the solution increases. For acetic acid, 0.05 M solution is % ionized whereas a 0.15 M solution is 1.0 % ionized. Chapter 16-Acids and Bases

18 Chapter 16-Acids and Bases
Polyprotic Acids Polyprotic Acids have more than one ionizable H+ The protons are removed in steps not all at once: It is always easier to remove the first proton in a polyprotic acid than the second. Therefore Ka1 > Ka2 > Ka3 etc. Chapter 16-Acids and Bases

19 Chapter 16-Acids and Bases
Weak Bases Weak bases remove protons from substances. There is an equilibrium between the base and the resulting ions: Example: The base dissociation constant, Kb, is defined as Chapter 16-Acids and Bases

20 Chapter 16-Acids and Bases
Weak Bases Trends in Weak Bases Bases generally have lone pairs or negative charges in order to attack protons. Most neutral weak bases contain nitrogen. Amines are related to ammonia and have one or more N-H bonds replaced with N-C bonds (e.g., CH3NH2 is methylamine). Anions of weak acids are also weak bases. Example: OCl- is the conjugate base of HOCl (weak acid): Chapter 16-Acids and Bases

21 Chapter 16-Acids and Bases
Some Weak Bases Chapter 16-Acids and Bases

22 Conjugate Acid/Base Strengths
When two reactions are added to give a third, the equilibrium constant for the third reaction is the product of the equilibrium constants for the first two: For a conjugate acid-base pair: The larger the Ka, the smaller the Kb. That is, the stronger the acid, the weaker the conjugate base. Taking negative logarithms: Chapter 16-Acids and Bases

23 Acid-Base Properties of Salt Solutions
Ionic salts dissociate completely Acid-base properties of salts are a consequence of the reaction of their ions in solution. The reaction in which ions produce H+ or OH- in water is called hydrolysis. Anions from strong acids are neutral. Anions from weak acids are basic: Polyatomic cations with ionizable protons can be considered conjugate acids of weak bases: Chapter 16-Acids and Bases

24 Chapter 16-Acids and Bases
Binary Acid Strengths Consider H-X: For this substance to be an acid we need: H-X bond to be polar H-X bond must be weak enough to be broken The conjugate base, X-, must be stable HF is a weak acid because the bond energy is high and not easily broken. H-I is a strong acid, easy to break H-I bond. Acid strength increases across a period and down a group. Base strength decreases across a period and down a group. Chapter 16-Acids and Bases

25 Acid-Base Behavior and Chemical Structure
Binary Acids Chapter 16-Acids and Bases

26 Chapter 16-Acids and Bases
Oxy-acids Oxyacids contain O-H bonds All oxyacids have the general structure Y-O-H. The strength of the acid depends on Y and the relative location of the electron density near Y: If Y is a metal (low electronegativity), then the substance is a base. If Y has intermediate electronegativity (e.g. I, EN = 2.5), the substance is a weak oxyacid. If Y has a large electronegativity (e.g. Cl, EN = 3.0), the electrons are located closer to Y than O and the O-H bond is polarized to lose H+. More O atoms attached to Y increase the O-H bond polarity and the strength of the acid increases (e.g. HOCl is a weaker acid than HClO2 which is weaker than HClO3 which is weaker than HClO4 which is a strong acid). Chapter 16-Acids and Bases

27 Acid-Base Behavior and Chemical Structure
Oxyacids Chapter 16-Acids and Bases

28 Acid-Base Behavior and Chemical Structure
Carboxylic Acids Carboxylic acids all contain the COOH group. All carboxylic acids are weak acids. When the carboxylic acid loses a proton, it generate the carboxylate anion, COO-. Chapter 16-Acids and Bases

29 Chapter 16-Acids and Bases
Lewis Acids and Bases Recall: Brønsted-Lowry acid is a proton donor. Focusing on electrons: a Brønsted-Lowry acid can be considered as an electron pair acceptor. Lewis acid: electron pair acceptor. Lewis base: electron pair donor. Note: Lewis acids and bases do not need to contain protons. Lewis acids generally have an incomplete octet (e.g. BF3). Transition metal ions are generally Lewis acids. Lewis acids must have a vacant orbital (into which the electron pairs can be donated). Chapter 16-Acids and Bases

30 Chapter 16-Acids and Bases
Metals and pH Hydrolysis of Metal Ions (can be tricky) Alkali and Alkaline-earth metals form bases in solution. Transition metal ions are positively charged and attract water molecules (via the lone pairs on O). The higher the charge, the smaller the metal ion and the stronger the M-OH2 interaction. Hydrated metal ions (water-complexes) act as weak acids: [H3O+] increases as the size of the ion decreases (Ca2+ / Zn2+) and as the charge increases (Na+ / Ca2+ and Zn2+ / Al3+). Na+ = no acidic behavior, Al3+ = very acidic cation Chapter 16-Acids and Bases

31 Chapter 16-Acids and Bases
Lewis Acids and Bases Hydrolysis of Metal Ions End of chapter 16 (phew) Chapter 16-Acids and Bases

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